1. Covalent Bonding & Lewis Structures

2. Types of Bonds- 3 Types Ionic (metal/nonmetal)- electron is transferred from the metal to the nonmetal Metallic (metal/metal)- electrons are said to be in an “electron sea”, or delocalized. They can move around between all the atoms. This explains why metals conduct electricity. Covalent (nonmetal/nonmetal)- electrons are shared so that all atoms end up with eight outer electrons

3. Covalent Bonding Covalent bonds involve the SHARING of electrons between atoms. Atoms are joined by a covalent bond, made up of two electrons and represented by a dash. Electrons are either shared in a bond, represented by a dash, or are lone electrons, represented by dots.

4. Representative Particles Covalent Compounds consist of individual particles called molecules. Molecules can have many different shapes depending on their atoms and electronegativites. Water molecules shown here have a bent shape.

5. Formation of a Hydrogen Molecule When a hydrogen atom encounters another hydrogen atom, they form a hydrogen molecule. Since they are both nonmetals, this will be a covalent bond. Each H atom has one electron to share. They form a single bond. A bond is when two electrons are shared.

6. Diatomic molecules Certain elements do not exist alone, because they are more stable when they are paired with another atom. Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 are the atoms that must either be written with a subscript of 2 or paired with another atom. The way we can remember which atoms are diatomic is to remember the goofy word Brinclhof.

7. Determining the number of bonds N – A = # of bonds 2 N is the # of needed electrons, which is 8 for all elements but H, which is 2 A is the # of available electrons, which is the number of valence electrons

8. 1) Determine the # of bonds needed using the given formula. 2) Draw a skeleton structure. “Central” means in the middle. The atom with the smallest subscript is usually the central one. 3) Connect the atoms with the # of bonds determined in step 1. 4) Finish by making sure all atoms in the structure have an octet of electrons. 5) Remember, if the structure has a positive charge, it has fewer available electrons and if it has a negative charge, it has more available electrons.

9. Carbon bonding Carbon follows the rules we have talked about in class, but it has a few unique properties that will help you draw structures for it.

10. Special rules for carbon 1) Carbon is always the central atom in its compounds and it tends to undergo bonding to itself. This tendency of carbon to bond to itself is so important that it has a special term to describe it – catenation. 2) Carbon always has four bonds. 3) Any time you have a choice for a skeleton structure, always go for the most symmetrical option. This is true for any compound, not just carbon’s.

11. Resonance Before we talk about resonance, there are a few concepts we should talk about. Bond length – A general principal is the more bonds between two atoms, the shorter the bond length will be. In other words, a single bond will be the longest, and a triple bond will be the shortest Bond strength – The more bonds between 2 atoms, the stronger that bond will be, and the more energy it will take to break that bond.

12. Now on to resonance. Resonance occurs when you have a double bond which can be placed in more than one location. A good example of resonance occurs in the carbonate polyatomic, CO 3 2-.

13. Exceptions to the octet rule All elements want an octet of electrons around them, except hydrogen, who wants two. Actually, there are two more elements which do not want an octet of electrons and those are Be (wants 4) and B (wants 6). Be careful and note that when you do a needed calculation for those elements the number is not 8.